A  Study  of  the  System  Ammonia- 
Water  as  a  Basis  for  a  Theory 
of  the  Solution  of  Gases 
in  Liquids 


EXCHANGE 

JIM  24  192, 


DISSERTATION 


SUBMITTED  TO  THE  BOARD  OF  UNIVERSITY  STUDIES  OF 

THE  JOHN  HOPKINS  UNIVERSITY  IN  CONFORMITY 

WITH  REQUIREMENTS  FOR  THE  DEGREE  OF 

DOCTOR  OF  PHILOSOPHY 


BY 


BENJAMIN  SIMON  NEUHAUSEN 


BALTIMORE 
June,  1921 


A  Study  of  the  System  Ammonia- 
Water  as  a  Basis  for  a  Theory 
of  the  Solution  of  Gases 
in  Liquids 


DISSERTATION 


SUBMITTED  TO  THE  BOARD  OF  UNIVERSITY  STUDIES  OF 

THE  JOHN  HOPKINS  UNIVERSITY  IN  CONFORMITY 

WITH  REQUIREMENTS  FOR  THE  DEGREE  OF 

DOCTOR  OF  PHILOSOPHY 


BY 

BENJAMIN  SIMON  NEUHAUSEN 


BALTIMORE 
June,  1921 


Acknowledgment 4 

Introduction 5 

The  Investigation  of  the  Vapor  Phase 

Apparatus 7 

Procedure 14 

Results 15 

The  Investigation  of  the  Liquid  Phase 

Apparatus 18 

Procedure 18 

Results 20 

Discussion  of  Results 24 

A  Theory  of  the  Solution  of  Gases  in  Liquids 24 

Summary 31 

Biography 33 


•t/l  QOCiO 


:  V 


ACKNOWLEDGMENT 

This  investigation  was  undertaken  at  the  suggestion  of 
Dr.  Walter  A.  Patrick,  to  whom  the  writer  wishes  to  express 
his  gratitude  for  very  helpful  guidance,  encouragement,  and 
kindness. 

The  writer  desires  to  thank  Drs.  Frazer,  Reid,  Lovelace, 
and  Thornton  for  laboratory  and  class  room  instruction  and 
for  helpful  suggestions  in  the  pursuit  of  this  investigation. 


A  STUDY  OF  THE  SYSTEM  AMMONIA-WATER  AS  A 
BASIS  FOR  A  THEORY  OF  THE  SOLUTION  OF  GASES 

IN  LIQUIDS 

Introduction 

To  a  great  extent  the  aim  of  most  investigations  on  the 
absorption  of  gases  in  water  and  other  liquids  has  been  to  note 
how  far  the  particular  gas  conformed  with,  or  deviated  from 
Henry's  law.  The  investigation  of  the  absorption  of  gases 
by  liquids  may  be  said  to  begin  with  Henry's1  study  of  the 
absorption  of  carbon  dioxide  by  water,  a  result  of  which  was  his 
famous  law  that  the  quantity  of  a  gas  absorbed  in  a  liquid  is 
proportional  to  the  pressure  of  the  gas.  Dalton's  theoretical 
explanation  of  this  law  on  the  basis  of  Boyle's  law,  and  his 
equally  well  known  law  of  partial  pressures  of  gases2  tended  to 
give  a  very  definite  trend  to  the  problem  of  absorption  of  gases 
by  liquids. 

Through  the  investigations  of  Bunsen3  on  the  absorption 
of  the  fixed  gases,  of  Schonfeld4  on  hydrogen  sulphide,  of 
Roscoe  and  Dittmar5  on  ammonia  and  hydrochloric  acid,  and 
of  Sims6  on  ammonia  and  sulphur  dioxide,  it  became  evident 
that  Bunsen  was  correct  in  his  stand  that  the  general  law  that 
expresses  the  relation  between  pressure  and  the  quantity  of  gas 
absorbed  is  a  complex  function  and  that  conformity  to  Henry's 
law  in  case  of  the  relatively  soluble  gases  was  the  exception 
rather  than  the  rule.  It  was  also  recognized  that  the  chemical 
properties  of  the  molecules  involved,  as  for  instance  their 
attraction  for  one  another  influenced  markedly  the  absorption. 


1  Phil.  Trans.,  1803;    Gilb.  Ann.,  20,  147  (1805). 

2  Gilb.  Ann.,  12,  385  (1803). 

3  Liebig's  Ann.,  93,  1  (1855). 

4  Ibid.,  95,  1  (1855). 

•  Ibid.,  112,  327  (1859). 
•Ibid.,    118,   333    (1861). 


*        *>«        a     ^*  *"•••.*  . 

•  •     : 
«  i.        •  •          •     •  • 


Thus,  for  example,  Roscoe  and  Dittmar  differentiate  clearly 
in  their  investigation  between  the  absorption  of  hydrochloric 
acid  gas  and  ammonia,  since  a  current  of  air  passed  through  an 
aqueous  solution  of  hydrochloric  acid  can  only  remove  gas 
up  to  a  definite  concentration  while  all  the  ammonia  may  be 
removed  from  an  ammonia  solution.  This  tendency  to  account 
for  absorption  of  gases  on  the  basis  of  physical  and  chemical 
properties  had  its  first  advocate  in  Th.  Graham1  who  in  a 
very  remarkable  paper  advanced  the  theory  that  gases  were 
actually  condensed  when  they  were  absorbed  by  a  liquid, 
and  that  solutions  of  gases  should  be  considered  as  binary 
liquid  mixtures  of  the  solvent  and  liquefied  gas.  To  this  view 
we  shall  return  in  the  discussion  of  the  results. 

The  interest  as  to  the  validity  of  Henry's  law  did  not  cease, 
and  a  number  of  investigations  on  the  solubility  of  carbon 
dioxide  were  pursued.  While  Khanikoff  and  Louguinine2 
found  that  Henry's  law  held  up  to  four  atmospheres,  v.  Wrob- 
lewski,3  Cassuto,4  and  Sander5  found  that  there  was  a  marked 
deviation  up  to  100  C. 

Ammonia  was  chosen  for  the  present  investigation  for  a 
number  of  reasons.  First,  it  has  a  very  high  solubility  in  water 
that  changes  very  markedly  with  the  pressure,  and  thus  there 
are  avoided  those  slight  absolute  changes  in  solubility  in  the 
case  of  a  gas  as  carbon  dioxide,  in  the  interpretation  of  which 
free  rein  is  given  to  preconceived  notions.  Secondly,  because 
of  its  economic  importance,  the  physical  constants  of  am- 
monia, such  as  vapor  tensions,  density  of  the  liquid,  surface 
tension,  etc.,  have  been  determined  more  accurately  than 
most  other  gases.  Thirdly,  it  has  a  relatively  high  critical 
temperature. 

The  first  investigator  to  work  extensively  on  the  absorp- 


1  Annals  of  Philosophy,  12,  69-74  (1826) ;    Chemical  and  Physical  Re- 
searches (Edinburgh,  1876),  pp.  1-6. 

2  Ann.  Chim.  Phys.,  (4),  11,  412  (1867). 

3  Wied.  Ann.,  17,  103  (1882);  18,  290  (1883). 

4  Nuovo  Cimento,  (5),  6,  (1903). 
6Zeit.  phys.  Chem.,  78,  513  (1912). 


tion  of  ammonia  was  Carius.1  His  methods  and  results  were 
proved  unreliable  by  Roscoe  and  Dittmar2  who  investigated 
the  absorption  of  ammonia  by  water  up  to  a  pressure  of  2000 
millimetres,  at  0°  C.  Sims3  checked  their  values  and  in- 
vestigated the  absorption  also  at  20°  C  and  40°  C  up  to  2100 
millimetres,  and  at  100°  C  up  to  1400  millimetres.  Watts4 
also  obtained  results  in  agreement  with  Sims,  while  Raoult5 
who  worked  at  one  atmosphere  and  varying  temperatures 
obtained  somewhat  higher  values  than  the  other  investigators. 
Perman6  who  worked  at  0°  C,  20°  C,  and  40°  C,  with  solu- 
tions up  to  22.5%  obtained  vapor  pressures  that  tally  fairly 
well  with  those  of  Sims.  Mallet7  working  at  a  pressure  of 
743-744.5  millimetres  obtained  the  solubilities  at  —10°  C, 
-20°  C,  -30°  C,  and  -40°  C. 

I.    The  Investigation  of  the  Vapor  Phase. 
1.    Apparatus 

The  composition  of  the  vapor  phase  of  binary  liquid 
mixtures  has  been  studied  both  by  dynamic  and  static  methods. 
The  principal  dynamic  method  was  to  allow  the  liquid  mixture 
to  boil  at  certain  pressures  and  temperatures  and  to  analyze 
the  distillate  and  residue  by  a  refractometer,  as  for  example, 
in  the  classical  researches  of  Zawidzki;8  or  to  bubble  air  through 
the  solution  and  absorb  the  vapors  in  suitable  solutions,  and 
thus  determine  the  composition  of  the  vapors,  as  was  done, 
for  instance,  by  Perman9  in  his  research  on  the  vapor  pressures 
of  ammonia  solutions.  Both  of  these  dynamic  methods  are 
open  to  rather  grave  objections.  First,  in  order  to  have  the 
vapor,  collected  as  distillate  or  in  suitable  solutions,  correspond 
with  any  degree  of  accuracy  to  the  solution  of  the  composition 

1  Liebig's  Ann.,  99,  129  (1856). 

2  Loc.  cit. 

3  Loc.  cit. 

4  Liebig's  Ann.  Suppl.,  3,  227   (1864). 

5  Ann.  Chim.  Phys.,  (5),  1,  262  (1874). 

6  Jour.  Chem.  Soc.,  79,  718  (1901);    83,  1168  (1903). 

7  Am.  Chem.  Jour.,  19,  804  (1897). 

8  Zeit.  phys.  Chem.,  35,  129  (1900). 

9  Jour.    Chem.    Soc.,    83,    1168    (1903). 


at  the  beginning  or  the  end  of  the  experiment,  one  has  to 
begin  with  a  large  volume  of  liquid;  and  moreover  one  must 
distill  off  only  a  small  portion.  Especially  in  cases  in  which 
one  of  the  components  has  a  high  vapor  pressure,  the  process 
should  be  continued  for  only  a  short  time.  It  is  very  question- 
able, as  Roozeboom1  observes,  whether  in  mixtures,  in  which 
there  is  a  great  difference  in  the  composition  of  both  phases, 
whether  the  vapor  that  is  formed  at  the  first  few  minutes 
corresponds  to  the  true  equilibrium.  In  fact,  impelled  by  such 
considerations,  Cunaeus2  preferred  to  let  the  vapor  remain 
in  contact  with  the  liquid  and  analyzed  the  same  by  optical 
methods.  Moreover,  as  in  the  present  investigation,  solutions 
of  high  concentrations  of  ammonia  were  used,  and  the  total 
vapor  pressure  of  the  solution  is  most  often  more  than  one 
atmosphere  even  at  0°  C,  distillation  consequently  was  made 
rather  difficult.  As  for  the  bubbling  method,  its  disadvan- 
tages have  already  been  expressed  by  Perman:3  "Unfor- 
tunately, the  method  is  not  applicable  over  a  very  wide  range 
of  temperature  or  with  very  varying  concentrations  of  the 
solution,  for  when  the  vapor  pressure  becomes  nearly  equal 
to  atmospheric  pressure,  a  very  little  air  will  draw  off  a  large 
quantity  of  vapor,  and  moreover  evaporation  of  ammonia 
becomes  so  rapid,  that  it  is  impossible  to  keep  the  temperature 
constant."  Another  obstacle  to  the  use  of  any  dynamic  method 
for  this  investigation,  can  be  seen  from  one  of  the  results 
obtained  at  0°  C.  An  ammonia  solution  which  has  a  partial 
pressure  of  1868  mm  ammonia  at  0°  C,  has  a  partial  pressure 
of  0.51  mm  of  water.  Now  assuming  ammonia  to  obey 
the  ideal  gas  laws,  to  obtain  even  5  mg  of  water,  22.8  litres 
of  vapor  reduced  to  standard  conditions,  or  practically  one 
gram  mole  of  ammonia  would  have  to  be  distilled  off,  a  volume 
too  large  to  deal  with,  and  moreover,  a  huge  quantity  of  solution 
would  be  required  in  order  to  be  able  to  neglect  the  loss  of  that 
much  ammonia. 


1  Heterogene  Gleichgewichte,  II,  part  I,  page  20. 

2  Zeit.  phys.  Chem.,  36,  232  (1901). 

3  Loc.  cit. 


9 

A  static  method  for  measuring  partial  pressures  was 
therefore  developed  on  the  basis  of  the  following  considerations. 

At  a  fixed  temperature,  a  solution  of  a  certain  definite 
composition  of  A  and  B,  has  a  fixed  total  pressure,  made  up 
of  two  well  defined  partial  pressures  of  the  vapors  of  A  and  B. 
B  is  assumed  to  have  a  greater  vapor  pressure  than  A.  Now, 
suppose  that  by  some  means  a  quantity  of  vapor  of  A  is 
forced  into  the  vapor  phase  of  the  system.  Since  the  partial 
pressure  for  A  is  exceeded,  some  of  it  will  condense.  As  some 
of  the  vapor  of  B  will  also  dissolve  in  this  condensate,  a  re- 
adjustment will  take  place,  and  a  solution  will  be  obtained 
which  is  somewhat  less  concentrated  with  respect  to  B,  and 
since  B  has  been  assumed  to  have  a  much  greater  pressure  than  A 
(in  the  present  case,  on  the  order  of  the  relative  vapor  pressures 
of  ammonia  and  water),  the  partial  pressure  of  B  is  appreciably 
diminished,  while  that  of  A  is  slightly  increased.  The  system 
therefore  will  suffer  a  lowering  in  the  total  pressure,  since  the 
volume  is  not  changed  appreciably  by  the  minute  quantity  of 
condensate. 

This  idea  was  then  extended  as  follows :  If  to  ammonia 
gas  at  a  certain  temperature  and  pressure  below  that  corre- 
sponding to  the  vapor  tension  of  liquid  ammonia  at  that  tem- 
perature, there  is  gradually  added  some  water  vapor,  the  pres- 
sure of  the  mixture  will  rise  until  the  partial  pressure  of  the 
water  vapor  equals  that  corresponding  to  the  partial  pressure 
of  water  vapor  over  a  solution  whose  partial  pressure  of  am- 
monia at  that  temperature  is  such  as  at  the  start.  Any  further 
addition  of  water  vapor  will  entail  condensation  of  water  with 
consequent  solution  of  some  ammonia  and  reduction  in  the 
total  pressure. 

In  Fig.  1  is  given  a  sketch  of  the  apparatus  employed. 
By  opening  reducing  valves  RI,  R2,  R3,  and  R4,  ammonia 
passed  in  from  the  ammonia  tank  into  the  manometer  and 
small  wrought  iron  tanks  Tj  and  T2.  The  auxiliary  manometer 
(A  M)  was  bent  at  an  obtuse  angle  of  157  degrees  so  that 
angle  "a"  had  a  sine  of  .1994.  Thus  a  rise  or  drop  of  one 


10 


Fig.  1 

millimetre  in  the  vertical  would  be  indicated  by  a  rise  or  drop 
of  five  millimetres  of  the  mercury  in  the  oblique  tubing  and 
could  be  read  on  the  attached  scale  "S"  to  within  0.2  mm. 
In  other  words,  change  of  pressure  in  tank  TI,  from  that  of 
tank  T2  could  be  measured  to  0.08  mm.  When  we  consider 
that  Perman1  only  claimed  an  accuracy  of  0.5  mm  for  his 
measurements  of  the  partial  pressure  of  water  vapor  in  total 
pressures  up  to  600  mm,  the  advantage  of  this  static  method 
can  be  seen. 

The  water  vapor  was  injected  by  means  of  the  water 
vapor  injector  (W  V  I)  which  was  connected  by  means  of 
valve  Rr,  to  1\.  The  distilled  water  contained  in  the  reservoir 
bulb  (W  R)  was  allowed  to  distill  over  into  W  V  I  which 
was  evacuated  in  the  following  manner.  Valves  R2  and  R3  were 
closed  to  prevent  any  direct  communication  between  TI  and  T2, 
while  R^,  was  opened.  Exhaust  valves  R6  and  R7  were  con- 
nected by  a  "Y"  tube  to  a  vacuum  pump  and  evacuated. 
That  was  continued  intermittently  for  several  hours  while 

1  Loc.  cit. 


11 


Si  was  shut  and  opened  at  intervals.  R5  and  Si  were  then  shut 
off,  and  ammonia  was  passed  into  Tj.  and  T2  through  R2  and  R3 
and  these  tanks  were  then  evacuated.  This  process  was  re- 
peated a  number  of  times  in  order  to  remove  any  water  vapor 
remaining  in  1\.  By  opening  R5  and  raising  the  mercury  level 
bulb  (M  L)  the  vapor  could  be  forced  into  IV 

M  T  was  a  movable  thermostat  that  could  be  raised 
and  lowered  by  means  of  pulley  and  tackle.  The  temperature 
was  regulated  to  0.05°  C.  It  was  necessary  to  have  this  mov- 
able thermostat  in  order  to  be  able  to  remove  it,  when  the 
ammonia  in  TI  and  T2  had  to  be  cooled  by  means  of  a  mixture 
of  carbon  dioxide  snow  and  ether  before  the  water  vapor  could 
be  forced  in. 

All  the  tubing  used  in  the  apparatus  was  made  of  brass. 
The  metal-glass  joint  between  the  water  vapor  injector  and 
R5  is  described  further  in  this  account. 

The  pressure  was  measured  by  means  of  a  continuous 
open  mercury  manometer,  which  consisted  of  four  glass  U- 
tubes,  each  having  a  length  of  1600  mm.  These  were  of 
hard  glass  with  a  6-mm  bore  and  a  wall  thickness  of  1  mm. 
It  is  diagrammatically  represented  in  Fig.  2.  The  structure 


A  a 


H 


Fig.  2 


into  which  the  glass  U-tubes  were  mounted  was  made  of  brass 
tubing,  and  the  glass  tubes  were  connected  to  it  by  means 
of  glass-metal  joints  (G  M  J)  described  below.  The  U- 
tubes  were  filled  half  their  length  with  redistilled  mercury, 
and  all  the  space  from  the  top  of  the  mercury  in  "B"  to  the 
mercury  in  "G,"  including  the  metal  structure,  was  filled 
with  benzene,  which  was  used  as  the  communicating  liquid. 
"H"  communicated  with  the  air  through  a  tube  filled  with 
soda  lime.  Since  each  tube  was  1600  mm  in  length,  a  total 
pressure  of  8  atmospheres,  after  making  various  corrections, 
plus  the  corrected  barometric  reading  at  the  time,  could 
be  measured,  or  approximately  9  atmospheres.  Readings 
were  made  to  the  nearest  millimetre,  so  that  the  maximum  error 
in  reading,  when  all  the  U-tubes  were  in  series,  could  be 
about  4  mm,  and  there  was  correspondingly  less  error  as 
fewer  tubes  were  read.  A  sliding  rod  "S  R"  that  could  slide 
over  two  fixed  posts  was  used  to  read  differences  in  level 
between  the  mercury  columns.  As  can  be  seen  from  the 
diagram,  there  was  no  need  for  correction  for  meniscus  de- 
pression. Valves  Vi  and  V2  were  always  open.  By  keeping 
valves  V3  and  V4  closed  the  pressure  was  measured  by  the 
whole  series.  By  opening  V3  the  second  U-tube  was  by-passed 
and  by  opening  both  V3  and  V4  both  the  second  and  third 
U-tubes  were  by-passed.  Vf)  was  used  as  an  exhaust  valve. 

The  idea  of  a  continuous  manometer  was  first  applied  by 
Richards  in  1845.  Thiesen1  and  then  Wiebe2  both  discussed 
and  described  such  manometers.  By  using  compressed 
gas  as  the  communicating  fluid,  H.  Kamerlingh  Onnes3  had 
found  such  an  arrangement  very  satisfactory  in  making  high 
pressure  measurements,  in  which  greater  accuracy  was  de- 
sired than  is  obtainable  with  a  closed  gas  manometer.  Re- 
cently at  the  Bureau  of  Standards,4  such  a  manometer  was 
used  to  measure  pressures  up  to  15  atmospheres  with  success, 

1  Zeit.  Instrumentenkunde,   1881,  p.   114. 

2Zeit.  comprimierte  Gase,  1897 -1898,  pp.  8,  25,  81,  107. 

3  Communications  Phys.  Lab.  Leyden,  Nos.  44,  67,  70  and  146. 

4  Jour.  Am.  Chem.  Soc.,  42,  206  (1920). 


13 


alcohol  being  used  as  the  communicating  liquid.  The  im- 
provement made  in  the  manometer  used  in  the  present  work 
was  that  the  U-tubes  were  all  of  practically  like  bore,  and 
so  pressure  could  be  read  at  any  point  desired,  while  in  those 
used  at  Leyden  and  at  the  Bureau  of  Standards,  the  tubes 
were  of  very  small  bore  with  the  exception  of  a  small  section 
in  the  lower  part  of  the  left  arm  and  in  the  upper  part  of  the 
right  arm.  Because  of  this  restriction  there  was  much  addi- 
tional manipulation  and  balancing  necessary,  and  pressures 
less  than  a  given  pressure  could  not  be  measured.  While 
the  arrangement  used  in  this  investigation  required  a  greater 
quantity  of  mercury,  yet  this  disadvantage  was  thought  to  be 
overbalanced  by  the  considerations  noted  above. 

Fig.  3  illustrates  the  method  employed 
in  making  glass-metal  joints.  The  tubing 
which  it  was  desired  to  connect  had  a  flange 
in  it  (F).  There  was  a  leather  gasket  on 
this  flange  (Li) .  The  glass  tubing  was  flared 
and  inserted  in  the  metal  collar  (M  C) 
and  made  fast  with  plaster  of  Paris,  (P 
P).  The  glass  surface  was  then  polished 
in  turn  with  rough  and  fine  emery  and 
rouge.  A  second  leather  gasket  (L2)  was 
put  between  the  collar  and  the  flange,  and 
the  two  were  put  firmly  together  by  means  of  the  nut  (N) 
that  slipped  over  the  flange  and  screwed  over  the  collar. 

The  ammonia  used  in  this  work  was  obtained  from  the 
Goetz  Ice  Machine  Co.,  Philadelphia,  Pa.,  and  was  kept  over 
metallic  sodium  to  remove  the  water.  The  hydrogen  that  was 
generated  was  allowed  to  escape  by  allowing  a  quantity  of 
ammonia  to  escape  twice  daily  for  a  week,  until  absence  of 
fixed  gases  was  indicated  by  the  fact  that  no  gas  remained 
behind  after  passing  500  cc  of  the  gas  through  sulphuric 
acid.  The  ammonia  which  was  to  be  used  in  the  investigation 
of  the  vapor  phase  was  then  distilled  into  a  second  cylinder, 
the  first  and  last  fractions  being  discarded.  In  the  experiments 


14 


on  the  solubility  of  ammonia,  it  was  found  unnecessary  to 
make  this  distillation,  since  the  ammonia  so  dried,  showed 
no  trace  of  fixed  gases,  and  since  before  starting  a  run  a  large 
volume  of  ammonia  was  always  run  through  the  apparatus 
to  remove  any  air. 

The  temperature  of  0  °  C  was  obtained  by  mixing  washed 
ice  shavings  with  water;  and  those  of  20°  C  and  40°  C  by 
means  of  an  electric  thermostat. 

2.    Procedure 

In  the  investigation  of  the  vapor  phase,  the  following 
procedure  was  followed.  Valves  RI,  R2,  and  R3  were  opened, 
and  a  vacuum  pump  connected  with  R6  and  R7.  After  the 
system  had  been  exhausted,  R6  and  R7  were  shut  off,  and  R4 
was  gradually  opened  until  there  was  a  pressure  of  about  two 
atmospheres  of  ammonia  in  the  tanks.  Valve  R4  was  then 
shut  off  and  the  system  again  evacuated.  This  was  repeated 
from  six  to  ten  times  before  every  run.  Ammonia  was  then 
allowed  to  come  into  the  system.  After  standing  in  the 
thermostat  for  about  half  an  hour,  after  which  time  the  mano- 
meter did  not  indicate  any  change  in  the  pressure,  valves 
RI,  R2,  and  R3  were  shut  off.  The  continuous  manometer  as 
well  as  the  barometer  were  then  read  and  the  temperature 
of  the  room  was  then  noted.  The  thermostat  was  then  lowered, 
and  moved  from  under  the  tanks  TI  and  T2,  and  TI  and  T2 
were  immersed  in  a  mixture  of  carbon  dioxide  snow  and  ether, 
the  ammonia  in  the  tanks  being  cooled  and  the  pressure  re- 
duced to  70-80  mm.  The  mercury  level  bulb  was  then  raised 
and  R5  opened  and  a  quantity  of  water  vapor  forced  in.  R5 
was  then  closed  and  TI  and  T2  were  immersed  in  a  beaker  of 
water  of  60°  C  for  about  ten  minutes,  and  then  the  thermostat 
was  raised  in  place.  After  a  half  hour  when  the  auxiliary 
manometer  showed  no  further  change  in  the  drop  of  the 
mercury  on  the  side  connected  to  Tb  the  drop  was  noted.  This 
injection  of  the  water  was  repeated  from  five  to  six  times  until 
the  mercury  in  this  side  began  to  rise,  indicating  a  drop  in 


15 

pressure  in  T\.  The  average  between  the  last  reading  at 
which  an  increase  in  pressure  was  indicated  and  the  first  at 
which  a  decrease  was  noted,  was  considered  as  the  partial 
pressure  of  the  water  vapor  corresponding  to  that  pressure  of 
ammonia  at  the  temperature  of  the  thermostat.  The  data 
thus  obtained  were  recalculated  as  shown  in  the  following 
specimen  record  sheet: 

Temperature  of  Bath :  20  °  C. 

Height  of  mercury  in  B  1053  mm 

Height  of  mercury  in  A  213  mm  840  mm 

Height  of  mercury  in  D  1049  mm 

Height  of  mercury  in  C  205  mm  844  mm 

Height  of  mercury  in  F  1051  mm 

Height  of  mercury  in  E  213mm  838mm 

Height  of  mercury  in  H  1102mm 

Height  of  mercury  in  G  268  mm  834  mm 

Total  mercury  height  in  manometer  3356  mm 

Room  temperature  near  manometer  23.75  C 
Factor  to  correct  for  expansion  of  Hg  0.9956 
Therefore  true  height  mercury  column  3341.1  mm 

Corrected  barometric  reading  757.3  mm 

Therefore  total  corrected  pressure  of  mercury  4098.4  mm  (Ri) 
Length  of  benzene  column  in  "  B "  617  mm  (Li) 

Difference  in  length  between  benzene  in  C  and  D  939  mm  (L2) 
Difference  in  length  between  benzene  in  E  and  F  838  mm  (La) 
Length  of  benzene  column  in  G  1037mm  CU) 

Equivalence  factor  of  benzene  at  23.75  C  to  mer- 
cury at  0  C  is  0.0643  (c) 
(L2-f;L3  +  L4-Li)c   ^                                 159.1    mm     (R2) 
RI  minus  R2  net  ammonia  pressure                3939.3   mm 
Reading  on  differential  manometer  8.25   mm  (max.) 

Reading  on  differential  manometer    7.25   mm     (first 

drop) 

8  25  +  7  25  2 

— -5 — '• — -  X  r  =  3.1  mm  partial  pressure  of  water. 

2i  o 

3.    Results 

The  results  obtained  are  given  in  the  following  table: 


16 

TABLE  I 

Composition  of  the  Vapor  Phase  of  Aqueous  Ammonia   Solutions 

Temperature  0°  C 


Partial  pressure  of 

Partial  pressure  of 

ammonia  in  mm 

water  in  mm 

1062 

1.04 

1100 

1.00 

1334 

0.84 

1868 

0.51 

2078 

0.35 

Temperature  20°  C 


Partial  pressure  of 
ammonia  in  mm 

Partial  pressure  of 
water  in  mm 

1146 

8.6 

1288 

8.3 

1445 

7.4 

2113 

5.9 

2112 

6.0 

2647 

4.5 

2624 

4.6 

3563 

3.7 

3725 

3.4 

3942 

3.1 

Temperature  40°  C 


Partial  pressure  of 
ammonia  in  mm 

Partial  pressure  of 
water  in  mm 

1120 

32.1 

1389 

29.8 

1579 

26.5 

1582 

26.4 

1926 

22.4 

2132 

20.8 

2381 

19.3 

2546 

18.9 

2969 

18.3 

3053 

18.3 

3395 

17.8 

3957 

17.4 

3928 

17.4 

17 


J 

£ 
E 

K 

0 
nj 

X 

\ 

o 
^, 

£ 

o 
in 

lO 

\ 

u 

cr 

L 

\ 

5 

<*• 

\ 

V 

\ 

\ 

TO' 
IK 

-AL  PF 
10      151 

ESSUI 
)0      19 

IE-  m 
50     m 

m 
10      19 

\ 

)0 

Fig.  4 

Partial  Pressure  HjO  and  Total  Pressure 
NH3-H2O  System  at  0°  C 

In  Figs.  4  and  5  are  plotted  the  partial  pressures  of  water 
as  ordinates  and  the  total  pressures  of  ammonia  plus  water 
as  abscissae  for  0°  C,  20°  C,  and  40°  C. 


^^ 


Partial  Pressure  H2O  and  Total  Pressure 
NH3-H2O  System  at  20°  C  and  40°  C 


18 


II. 


Fig.  6 


Investigation  of  the  Liquid  Phase. 
1.    Apparatus 

In  Fig.  6  is  a  sketch  of  the  apparatus 
for  determining  the  solubility  of  ammonia 
in  water  at  different  pressures  at  0°  C,  20° 
C,  and  40°  C. 

T,  the  solubility  tank  of  about  500  cc 
in  capacity,  was  connected  to  the  ammonia 
tank  and  to  the  manometer  by  means  of 
the  metal  joint  (M  J),  which  as  can  be  seen 
from  the  figure  was  similar  in  construction 
to  the  glass-metal  joints.  When  it  was  de- 
sired to  pass  ammonia  in  or  to  measure  the 
pressure  of  the  gas  phase,  the  connection  by 
means  of  M  J  would  be  made,  and  valves 
Vi  and  V2  opened.  A  sample  of  the  solu- 
tion could  be  obtained  in  the  small  tank 
S  T  which  unscrewed  from  T  at  C.  An 
eccentric  was  used  to  shake  the  solution  by 
causing  the  tank  to  make  a  60-degree  arc  of 
a  circle  in  the  thermostat. 

The  temperature  was  regulated  to  0.05° 
C. 


The  pressure  was  measured  on  the  same  manometer  de- 
scribed above. 

2.    Procedure 

The  tank  T,  without  the  sample  tank  S  T,  was  filled 
with  water  and  connected  to  the  cylinder  of  purified  ammonia. 
By  opening  Vz  about  300  cc  of  water  was  displaced,  and  about 
200  cc  of  water  with  an  atmosphere  of  ammonia  remained. 
The  sample  tank  was  now  screwed  on  to  T,  V3,  V4,  and  V5 
remaining  shut  and  ammonia  was  allowed  to  pass  in  for  ten 
minutes.  The  tank  was  then  disconnected  and  shaken  for 
about  ten  minutes,  and  then  more  ammonia  was  passed  in. 
When,  after  this  procedure  had  been  repeated  for  a  number  of 
hours,  the  pressure  reading  reached  a  value  about  that  desired ; 


19 

no  more  ammonia  was  passed  in,  and  shaking  was  continued 
for  fifteen  minutes  at  intervals,  until  pressure  readings  changed 
by  only  several  millimetres,  and  equilibrium  was  then  con- 
sidered to  have  been  reached.  The  pressure  registered  by  the 
manometer,  the  barometer  reading,  and  the  temperature 
of  the  room  were  noted.  Vr,  was  then  connected  to  a  vacuum 
pump,  and  V4  and  Vr,  opened,  and  S  T  evacuated.  Vr,  was 
then  closed  and  Vs  opened.  The  pressure  above  faced  the 
solution  into  S  T.  After  about  a  minute  V3  and  V4  were 
shut  and  S  T  unscrewed. 

The  above  method  of  drawing  off  a  sample  was  found 
not  to  affect  the  concentration  of  the  solution  by  the  increase 
of  the  vapor  phase  during  the  short  time  that  the  sample 
was  withdrawn.  For  example,  a  second  sample  taken  ten 
minutes  later  weighed  only  8  mg  more,  and  on  analysis  con- 
tained 1/10  of  1%  less  ammonia  than  the  first  sample.  This 
method,  moreover,  had  the  advantage  that  equilibrium  did 
not  have  to  be  reached  through  a  small  opening  in  the  valve. 

The  sample  tank  was  carefully  dried  with  filter  paper  and 
.weighed.  The  solution  was  then  run  into  a  measured  quantity 
of  standard  sulphuric  acid,  the  sample  tank  being  thoroughly 
washed  out.  S  T  was  dried  by  compressed  air  and  weighed. 
Every  third  or  fourth  time,  it  was  filled  with  mercury  and 
weighed  to  determine  the  internal  volume,  so  as  to  ascertain 
whether  any  change  had  taken  place  because  of  corrosion. 

An  aliquot  part  of  the  sulphuric  acid  solution  was  titrated 
with  dilute  sodium  hydroxide  solution,  methyl  red  being  used 
as  the  indicator.  Thus  the  quantity  of  ammonia  present 
was  determined. 

A  sample  record  sheet  of  one  solubility  determination 
will  illustrate  the  procedure: 

Temperature  of  bath:     20°  C 

Height  of  mercury  in  B  727  mm 

Height  of  mercury  in  A  529  mm      198  mm 

Height  of  mercury  in  H  767  mm 

Height  of  mercury  in  G  571  mm     196  mm 


Total  mercury  height  in  manometer          394  mm 
Room  temperature  near  manometer          28  C 
Factor  to  correct  for  expansion  of  mercury    0.9948 
Therefore  true  height  of  mercury  column       392  mm 
Correct  barometric  reading  765  mm 


Total  corrected  mercury  height  1157  mm  (Ri) 

Difference  in  columns  of  benzene  between  G  and  B    156  mm  (L) 
Equivalence  factor  of  benzene  at  28  °  C  to  mercury 

atO°C  0.0638(c) 

L  X  c  =  R2  10  mm 

RI  —  R2  =  Total  vapor  pressure  of  solution  1147  mm 

Weight  of  sample  tank  and  solution  196.001  g 

Weight  of  sample  tank  186 . 195  g 


Weight  of  solution  9 . 806  g 

222.91  cc  1.0532  N  H2S04  neutralized  3.991  g  NH3 

Water  present  in  sample  5.815g 
Therefore  0.686  g  NH3  per  gram  water 

Weight  sample  tank  and  mercury  at  24°  C  339.835  g 

Weight  sample  tank  186 . 195  g 

Weight  of  mercury  at  24  °  C  J53 . 640  g 

Therefore  volume  1 1 . 35  cc 

Therefore  density  of  solution  0 . 864 

Now  by  calculating  from  Berthoud's1  values,  the  density 
of  liquid  ammonia  at  0°  C,  20°  C,  and  40°  C,  and  from  the 
density  of  water  at  these  temperatures,  there  could  be  obtained 
the  densities  that  the  ammonia  solutions  would  have  if  no 
contraction  in  volume  of  the  water  and  ammonia  had  taken 
place.  From  these  calculated  values  and  densities  found  for 
the  solutions,  the  contraction  in  volume  could  be  calculated. 
Also,  by  means  of  curves  in  Figs.  4  and  5,  the  partial  pressures 
of  water  vapor  in  the  total  pressures  observed  could  be  read  off. 

3.    Results 

In  Table  II  the  solubility  data  obtained  is  summarized. 

Because  of  the  very  large  difference  in  the  vapor  pressures 
of  ammonia  and  water,  it  was  impossible  to  plot  the  partial 
pressure  curves  of  both  the  ammonia  and  water  on  one  diagram. 
To  be  able  to  compare  the  relative  trend  of  the  two  curves 

1  Helv.  Chim.  Acta,  1,  84-87  (1918). 


30      NH, 
H20 

Fig.  7 

Reduced  Partial  Pressure  Curves  of 
NH3andH2OatO°C 


Fig.  8 

Reduced  Partial  Pressure  Curves  of 
NHjandH2Oat20°C 


V 


\ 


0 
100 


20 
80 


40 
60 


60 
40 


K-.0 


Fig.  9 

Reduced  Partial  Pressure  Curves  of 
NH3andH2Oat40°C 

in  one  diagram,  there  have  been  plotted  in  Figs.  7,  8,  and  9, 
the  partial  reduced  pressures  of  the  ammonia  and  water,  for 
the  various  mole  fractions  of  ammonia  at  0°  C,  20°  C,  and 
40°  C.  The  reduced  partial  pressures  are  the  ratios  between 
the  observed  partial  pressures  and  the  pressures  of  the  pure 
components  at  these  temperatures.  TTI  and  7r2  indicate  the 
ammonia  and  water  curves,  respectively.  Points  marked  with 
an  X  are  taken  from  Perman's  work. 


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24 

Discussion  of  Results 

From  Figs.  7,  8,  and  9,  it  is  noted  that  the  partial  pressure 
curves  of  the  ammonia  are  concave  to  the  straight  line  curve 
of  ideal  mixtures,  and  that  the  curves  tend  to  approach  this 
line  with  increasing  concentration  of  ammonia.  This  ap- 
proach becomes  more  marked  with  increasing  temperature. 
Whether  these  curves,  if  continued,  would  cut  the  straight  line 
curve,  and  then  become  concave  to  the  axis  of  abscissae  (i.  e., 
a  negative-positive  curve  as  in  the  pyridine-water  system  in- 
vestigated by  Zawidzki) l  or  whether  the  curve  would  approach 
the  straight  line  and  become  asymptotic  to  it  cannot  be  fore- 
seen. Yet  from  analogy  with  the  curves  of  the  partial  pressures 
of  water  it  is  very  probable  that  the  ammonia  curves  become 
asymptotic  to  the  straight  line;  or,  in  other  words,  that  in 
dilute  solutions  of  water  in  ammonia,  the  partial  pressures 
of  the  ammonia  would  approach  those  required  by  Raoult's 
Law. 

From  the  curves  of  the  partial  pressures  of  the  water, 
it  is  seen  that  the  straight  line  is  approached  as  the  mole 
fraction  of  the  water  increases.  As  will  be  noted  by  comparison 
with  Table  II,  the  regions,  in  which  flex  points  appear,  corres- 
pond to  those  concentrations  in  which  there  is  a  maximum  or 
minimum  contraction  in  volume.  Nothing  further  can  be 
advanced  at  present  to  account  for  this  phenomenon. 

Since  the  vapor  pressure  of  ammonia  is  so  much  greater 
than  that  of  water,  the  total  vapor  pressure  curve  will  be 
that  of  the  ammonia  curve,  and  therefore  the  ammonia  so- 
lution will  most  probably  not  show  a  maximum  or  minimum 
boiling  point,  as  had  already  been  stated  by  Konowalow.2 

A  Theory  of  the  Solution  of  Gases  in  Liquids 

The  fundamental  concept  of  the  theory  of  the  solution 
of  gases  in  liquids,  which  it  is  desired  to  advance,  is  that 
first  enunciated  by  Graham3  in  his  remarkable  paper,  namely 

1  Zeit.  phys.  Chem.,  55,  129  (1900). 

*  Ber.  deutsch  chem.  Ges.,  17,  1531  (1884). 

3  Loc.  cit. 


25 

that  "gases  owe  their  absorption  by  liquids  to  their  capability 
of  being  liquefied  and  to  the  affinities  of  liquids  to  which  they 
become  in  this  way  exposed,"  and  that  "solutions  of  gases  in 
liquids  are  mixtures  of  a  more  volatile  with  a  less  volatile 
liquid."  Since  the  condensation  of  the  gas  is  regarded  as  the 
fundamental  phenomenon  of  its  absorption  by  a  liquid,  and 
solutions  of  gases  in  liquids  are  liquid  mixtures,  the  following 
postulates  follow: 

1.  A  gas  cannot  be  dissolved  in  appreciable  quantities 
above  its  critical  temperature,  since  above  the  critical  tem- 
perature   liquefaction    is    impossible.     The    fixed    gases   are 
dissolved  to  some  extent  by  water  and  other  liquids,  but  if 
absorption  coefficients  are  discarded,  and  the  relative  number 
of  gas  molecules  dissolved  by  the  water  molecules  considered 
instead,   the  exceedingly  small  quantity  becomes  apparent. 
Thus  according  to  v.  Antropoff1  one  gram  of  water  absorbs 
0.123  cc  Xenon  reduced  to  N.  T.  P.,  Xenon  being  the  most 
soluble  of  the  noble  gases.     In  other  words,  0.055  of  a  gram 
mole  of  water  absorbs  0 . 0555  of  a  gram  mole  of  Xenon,  or  one 
molecule  of  Xenon  for  10,000  molecules  of  water.     It  is  not 
true  liquefaction  that  is  taking  place  in  such  cases,  but  rather 
a  retention  of  the  Xenon  molecule  by  the  attractive  forces 
of  numerous  water  molecules. 

2.  Since  solutions  of  gases  in  liquids  are  to  be  considered 
as  liquid  mixtures,  the  degree  of  miscibility  of  the  liquids 
are  of  great  importance.     Thus,  since  liquid  carbon  dioxide 
and  liquid  sulphur  dioxide  have  only  a  limited  miscibility  with 
water,  their  gases  should  dissolve  in  water  to  a  limited  extent, 
for  any  liquefied  gas  separating  out  could  not  remain  as  such 
except  at  a  partial  pressure  equal  to  its  vapor  tension  at  that 
temperature.     The  relative  size  and  shape  of  the  pores  in  the 
liquid  would  also  tend  to  influence  the  quantity  of  gas  dissolved. 
Thus  Just2  found  that  the  solubility  of  carbon  dioxide  in 
various  liquids  was  greater  the  smaller  the  refractive  index. 


1  Zeit.  Elektrochemie,  25,  269  (1919). 

2  Zeit.  phys.  Chem.,  37,  343  (1901). 


26 

This  phenomenon  could  be  explained  on  the  basis  of  the 
Clausius-Mossotti  formula  and  the  electro-magnetic  theory 
of  light,  whereby 


in  which  u  is  the  true  space  occupied  by  the  molecules  and 
n  is  the  refractive  index.  Likewise,  solubility  might  be  in 
some  way  related  to  a  quantity  such  as  "b"  of  van  der  Waals. 

3.  The  chemical  nature  of  the  gas  molecules,  such  as 
polarity  would  likewise  influence  the  degree  of  solution. 
This  was  already  recognized  by  Roscoe  and  Dittmar1  in  their 
investigation  of  the  solubilities  of  ammonia  and  hydrogen 
chloride  in  water.  Thus,  while  they  denied  Bineau's  con- 
tention that  the  constant  boiling  mixtures  of  hydrochloric 
acid  were  definite  compounds,  they  nevertheless  concluded 
that  between  water  and  hydrochloric  acid  there  is  an  attraction 
appreciably  different  from  that  with  other  gases.  The  other 
inorganic  acids  forming  constant  boiling  mixtures  are  also 
strongly  polar  substances. 

On  the  basis  of  the  above  considerations,  the  solution  of 
gases  in  liquids  may  be  classified,  as  follows: 

I.  Solution  of  gases  above  their  critical  temperature. 

II.  Solution  of  gases  below  their  critical  temperature. 

(1)  Gases,  whose  condensates  have  limited  miscibility 
with  water,   and  which  are  only  limitedly  soluble, 
e.  g.,  CO2  and  SO2. 

(2)  Gases,  whose  condensates  are  miscible  with  water 
in  all  proportions, 

(a)  Those  of  strong  polarity  which  form  constant 
boiling  mixtures,  as  HC1,  and  are  extremely 
soluble. 

(b}   Those  whose  polarity  is  not  as  strong,  as  NH3, 

which  can  be  driven  entirely  out  of  solution. 
Patrick  and  McGavack2  in  their  theoretical  discussion  of 


1 1/oc.  cit. 

*  Jour.  Am.  Chem.  Soc.,  42,  946  (1920). 


27 

the  results  obtained  in  their  investigation  of  the  adsorption 
of  sulphur  dioxide  by  silica  gel,  attacked  the  problem  of  ad- 
sorption in  the  following  manner.  Postulating  that  when 
any  gas  is  adsorbed  in  appreciable  quantity,  condensation  of 
the  gas  is  actually  taking  place,  the  question  arises  how  one 
can  account  for  the  fact  that  the  gas  pressures  in  equilibrium 
are  much  lower  than  the  vapor  tension  of  the  liquefied  gas. 
These  authors  on  the  basis  of  capillarity  advanced  the  thesis 
that  the  lowering  of  the  vapor  pressure  was  due  to  the  fact 
that  the  liquid  was  under  a  tension  or  a  negative  pressure. 
It  is  well  known  that  a  liquid  under  a  hydrostatic  pressure 
has  a  greater  vapor  pressure  than  when  under  the  pressure  of 
merely  its  own  vapor,  and  conversely  when  it  is  under  a  tension 
or  negative  pressure,  one  would  expect  a  lowering  in  the  vapor 
pressure.  This  tension  could  be  calculated  on  the  basis  of  the 
Gibbs'  relation 

V 


in  which  dp  equals  the  change  in  vapor  pressure,  dP  equals 
change  in  hydrostatic  pressure,  V  equals  volume  of  condensed 
liquid  phase,  and  v  equals  volume  of  the  gas.  Furthermore, 
this  tension  must  cause  a  dilation  of  the  liquefied  gas  to  an 
extent  that  is  proportional  to  the  compressibility  of  the  liquid. 
The  compressibility  of  the  liquid  was  taken  as  some  function 
of  the  surface  tension.  They  finally  developed  the  following 
formula  which  tallied  very  well  with  experimental  results: 

Per  \i 


in  wrhich  V  is  the  volume  of  the  condensed  gas  absorbed  per 
gram  of  gel;  i.  e.,  the  mass  of  the  gas  adsorbed  divided  by  the 
density  of  the  liquefied  gas  at  the  temperature;  P  is  the  equi- 
librium gas  pressure;  <7  is  the  surface  tension  and  P0  the  vapor 
tension  of  the  liquefied  gas  at  the  temperature.  K  and  l/n 
are  constants.  By  plotting  the  values  of  log  V  as  ordinates 

and  log  ——  —  as  abscissae,  the  experimental  values  of  Patrick 
"o 

and  McGavack  fell  on  a  straight  line. 


28 


^ 

. 

<^ 

1 

rX 

[** 

X 

LO 

ofc 

-0?     -C 

LI         *fl 

1        *< 

1       *« 

5       +0 

.7        »ll 

.9     *l 

1 

Fig.  10 

Solubility  of  NH3  in  H2O. 
O  at  20°  C; 


O  at  0' 
at  40°  C 


Since  the  solution  of  gases  in  liquids  is  also  considered 
as  a  condensation,  the  results  obtained  in  this  investigation 
were  also  plotted  according  to  this  formula  as  can  be  noted 
from  Fig.  10.  The  densities  and  values  of  the  surface  tension 

of  liquid  ammonia  at  0°  C, 
20°  C,  and  40°  C,  were 
interpolated  from  the  values 
of  Berthoud1  and  the  vapor 
tensions  used  were  those 
obtained  at  the  Bureau  of 
Standards. 2  The  results  ob- 
tained for  0°  C,  20°  C,  and 
40  °C,  with  solubilities  rang- 
ing from  0.315  gram  to 
1.885  grams  of  ammonia 
per  gram  of  water  are  so  well  represented  by  this  formula,  that 
it  may  be  assumed  that  a  phenomenon  similar  in  its  manifesta- 
tions to  that  of  the  adsorption  of  sulphur  dioxide  by  silica 
gel  is  occurring  in  the  case  of  the  solution  of  ammonia  in  water. 
In  Curve  I,  Fig.  11,  other  available  data  on  the  solubility 
of  ammonia  in  water  at  lower  pressures  obtained  by  Perman, 
Roscoe  and  Dittmar,  and  Sims,  at  0°  C,  20°  C,  and  40°  C 
and  by  Mallet3  for  pressures  of  743  to  744.5  mm,  at  —10° 
C,  -20°  C,  -30°  C,  and  -40°  C,  as  well  as  the  results 
obtained  in  this  investigation  were  plotted.  The  results 
obtained  in  this  investigation  as  well  as  by  other  investigators 
at  varied  temperatures  and  pressures,  and  concentrations  as 
high  as  2.746  grams  ammonia  per  gram  of  water  fall  on  the 
same  straight  line  curve.  This  is  apparently  a  general  law 
for  the  solution  of  ammonia  in  water. 

On  Curve  II,  same  figure,  have  been  plotted  in  like  manner 
the  solubility  of  hydrogen  chloride  in  water  as  found  by  Roscoe 
and  Dittmar4  at  0°  C,  and  pressures  varying  from  58  to  1270 

1  Helv.  Chim.  Acta,  I,  84-7  (1918);  Jour.  Chim.  phys.  16,429  (1918). 

2  Jour.  Am.  Chem.  Soc.,  42,  206  (1920). 

3  Am.  Chem.  Jour.,  19,  804  (1897). 

4  Loc.  cit 


29 


mm.  The  surface  tension  of  liquid  hydrogen  chloride  was 
calculated  from  the  work  of  Mclntosh  &  Steel1  to  be  7.6 
dynes;  the  density  of  liquid  hydrogen  chloride  was  taken  as 
0.908  according  to  Ansdall;2  and  vapor  tension  of  liquid  hydro- 
gen chloride  was  taken  as  19,900  mm  according  to  Faraday.3 


-1.6       -12      -.8       -A 


LOG$ 

0       +4 


»8      +1.2 


Fig.  11 

On  Curve  I  O  indicates  observations  at  0°C;  •  at  20°  C;  O 

at  40 °C.   t  appended  indicates  data  of  Sims;  t  data  of 

Perman;  and  R  data  of  Mallet.     On  Curve  III  O 

indicates  data  at  7°C  and  •  data  at  20 °C 

On  Curve  III,  there  has  been  plotted  the  solubility  of 
sulphur  dioxide  in  water,  as  found  by  Sims4  at  7°  C  from 
pressures  of  27  to  1291  mm,  and  at  20°  C,  from  32.4  to 
1911  mm.  The  values  of  the  surface  tension  of  liquid  sulphur 
dioxide  at  7°  C  and  20°  C  were  calculated  from  values  given 
by  Landolt-Bornstein  to  be  27.2  and  24.5  dynes,  respectively. 


1  Zeit.  phys.  Chem.,  55,  141  (1906). 

2  Proc.  Roy.  Soc.,  30,  117. 

3  Phil.  Trans.,  135,  I,  155  (1845). 

4  Loc.  cit. 


30 

The  density  of  liquid  sulphur  dioxide  was  interpolated  from 
the  data  of  Cailletet  and  Matthias1  and  taken  as  1.383  for 
20°  C,  and  1.42  for  7°  C.  Regnault's  values  for  the  vapor 
tensions  were  used.2 

On  Curve  IV,  was  plotted  the  solubility  of  carbon  dioxide 
in  water  at  760  mm  pressure  at  from  0°  C  to  25°  C.  The 
solubility  data  were  those  of  Bohr  and  Bock.3  The  values 
of  surface  tension  at  these  temperatures  were  obtained  by 
interpolation  from  data  given  by  Landolt-Bornstein,  the 
values  used  being  as  follows ; 

Temperature,  C  0  5  10  15  20  25 
Surf  ace  Tension  4.65  3.5  2.74  1.82  1.00  .50  dynes 
The  density  of  liquid  carbon  dioxide  at  these  temperatures 
was  obtained  from  the  data  of  Warburg  and  v.  Babo.4  The 
data  of  Th.  Tate5  on  the  vapor  tension  of  liquid  carbon  dioxide 
were  used. 

From  the  Curves  I  to  IV  it  is  evident  that  for  each 
individual  gas  the  law  contained  in  the  formula 


holds  very  well,  but  that  these  lines  do  not  coincide.  From 
a  consideration  of  the  nature  of  the  solution  of  gases  in  liquids 
such  as  outlined  above  no  such  coincidence  would  be  expected. 
If  this  formula  is  written  in  logarithmic  form,  there  is  obtained 
the  following  equation  : 


log  V  =  log  K  +   ~  log  — 
n          t*j 

In  the  above  equation  when  —  is  set  equal  to  1,  log  K  becomes 


1  Comptes  rendus,  104,  1565  (1887). 

2  Landolt-Bornstein  Tabellen. 
3Wied.  Ann.,  44,  318  (1891). 

4  Ber.  Berl.  Akad.,  p.  509  (1882). 
6  Phil.  Mag.,  (4),  26,  502. 


31 

equal  to  log  V.  l/n  is  obviously  the  slope  of  the  straight 
line..  If  solutions  of  gases  in  liquids  are  considered  as  binary 
mixtures,  it  is  evident  that  for  the  same  solvent,  those  gases 
which  in  the  liquefied  state  mix  in  all  proportions  with  water 
(an  indication  that  their  molecular  forms  are  such  that  the 
molecules  of  the  one  liquid  fit  into  the  pores  of  the  other 
liquid)  their  condensates  will  be  taken  up  in  greater  quantity 
than  those  of  gases  whose  condensates  have  only  a  limited 
miscibility  with  water.  Besides  the  degree  of  miscibility, 
such  factors  as  the  dielectric  constant  of  the  liquefied  gas 
(which  is  a  function  of  the  space  occupied  by  the  molecules), 
no  doubt  play  an  important  r61e  in  determining  the  degree  of 
solubility.  It  is  proposed  that  the  values  of  the  constants 
K  and  l/n  depend  on  such  factors  as  miscibility  and  dielectric 
constants  of  the  liquefied  gases,  etc.  In  fact  from  a  comparison 
with  Fig.  1 1  it  becomes  evident  that  K  is  much  greater  for  the 
solubility  curves  of  ammonia  and  hydrogen  chloride,  which 
in  the  liquid  state  are  miscible  with  water  in  all  proportions 
than  for  carbon  dioxide  or  sulphur  dioxide  whose  condensates 
have  only  a  limited  solubility  in  water.  Also,  in  the  case  of 
these  four  gases,  it  has  been  found  that  the  values  of  l/n  par- 
allel to  some  extent  the  values  of  the  dielectric  constant  of  the 
liquefied  gas  dissolved,  i.  e.,  those  gases  whose  liquids  possess  a 
high  dielectric  constant  have  a  large  value  of  l/n.  But 
unfortunately,  there  are  not  enough  data  available  for  other 
gases  to  test  out  the  validity  of  this  relation. 

However,  this  fact  is  worthy  of  note — that  the  values 
of  K  and  l/n  for  each  particular  gas  are  independent  of  the 
temperature  and  the  partial  pressure  of  the  dissolved  gas. 
This  fact  is  of  great  importance;  and,  though  only  speculation 
as  to  its  significance  is  possible  at  present,  it  offers  very  fertile 
fields  for  investigation  as  to  the  factors  on  which  the  values 
of  these  constants  depend  in  the  application  of  the  adsorption 
formula  to  the  case  of  the  solubility  of  gases  in  liquids. 

Summary 

(1)  A  static  method  has  been  developed  for  measuring 


32 

the  partial  pressure  of  a  component  which  is  relatively  very 
small  compared  to  the  partial  pressure  of  the  second  component. 

(2)  This  method  has  been  used  to  determine  the  partial 
pressures  of  water  and  ammonia  of  concentrated  ammonia 
solutions  at  0°  C,  20°  C,  and  40°    C,    at   partial   pressures 
of  ammonia  varying  from  1000  to  4000  mm.     The  partial 
pressures  of  the  ammonia  were  measured  to  within  4  to  2 
millimetres;  and  those  of  the  water  to  0.08  millimetre. 

(3)  The  solubility  of  ammonia  in  water  was  determined 
at  0°  C,  20°  C,  and  40°  C  at  pressures  from   750   to   3600 
mm.     The  densities  of  these  solutions  were  also  determined. 

(4)  A  theory  of  the  nature  of  solutions  of  gases  in  liquids 
first  advanced  by  Graham,  has  been  amplified,  and  solutions 
of  various  gases  in  liquids  classified  on  the  basis  of  some  of 
the  physical  and  chemical  properties  of  the  gas. 

(5)  The  formula 


has  been  found  to  represent  well  the  solubility  of  ammonia 
hydrogen  chloride,  sulphur  dioxide,  and  carbon  dioxide  in 
water  at  varied  temperatures  and  pressures.  In  this  formula 
V  is  the  volume  occupied  by  the  liquefied  gas  dissolved  per 
gram  of  water;  P0is  the  vapor  tension  and  <r  the  surface  tension 
of  the  liquefied  gas  at  the  temperature  while  P  is  the  equilibrium 
gas  pressure.  The  constant  K  for  ammonia  has  the  value 
0.49  and  l/n  has  the  value  0.69. 

(6)  Solubility  data  of  HC1,  SO2,  and  CO2  have  also  been 
plotted   according  to  this  formula. 


BIOGRAPHY 

Benjamin  Simon  Neuhausen  was  born  July  31,  1896. 
His  early  education  was  received  in  the  public  and  high  schools 
of  New  York  City.  In  1918  he  received  the  A.  B.  degree  from 
Johns  Hopkins  University.  During  the  year  1919-1920  he 
was  a  Hopkins  Scholar  and  Student  Assistant;  during  1920- 
1921  he  was  a  Du  Pont  Fellow. 


Gay  lord  Bros. 

MaUers 

Syracuse,  N.  Y. 
ta.  w  21,  taoa 


YD  04909 


543992 


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